Brønsted–Lowry theory provides a deeper understanding of acid-base chemistry. Through these notes, delve into the nuances of proton donors and acceptors, and trace the journey of how our knowledge of acids has developed over the centuries.
Introduction to Brønsted–Lowry Theory
In the realm of acid-base chemistry, several theories have been proposed to offer an understanding of these reactions. One such influential theory is the Brønsted–Lowry theory. Proposed by Johannes Brønsted and Thomas Lowry independently in 1923, this theory centres around protons.
- Brønsted–Lowry Acid: A molecule or ion that can donate a proton (H+) to another substance. This act of donation characterises its acidic nature.
- Brønsted–Lowry Base: A molecule or ion that can accept a proton (H+) from another substance. The act of accepting a proton is what gives it its basic nature.
This theory is more encompassing than earlier definitions as it is not restricted to aqueous reactions.
Image courtesy of SAMYA
Analysing Reactions to Identify Acids and Bases
Using the Brønsted–Lowry theory, one can dissect a reaction to identify which participant is the acid and which one is the base.
Take the reaction between hydrogen chloride and water:
HCl + H2O -> H3O+ + Cl-
In this equation:
- HCl donates a proton to water, so it is the Brønsted–Lowry acid.
- H₂O receives the proton from HCl, hence it's the Brønsted–Lowry base.
Key points to remember:
- The compound that gives away the proton in the reaction is the acid.
- The compound that takes on the proton is the base.
Base vs. Alkali: Clarifying the Terms
The distinction between "base" and "alkali" is often blurred in everyday language. However, in chemistry, precision is paramount.
- Base: Any substance, irrespective of its solubility, that can accept a proton is termed a base. This is a broad category that includes numerous substances.
- Alkali: An alkali, on the other hand, specifically refers to a base that dissolves in water. This solubility criterion narrows down the substances that can be termed alkalis. Common examples include sodium hydroxide (NaOH) and potassium hydroxide (KOH).
Image courtesy of Passnownow
Tracing the Evolution of Acid Definitions
The story of how we've come to understand acids is a testament to the ever-evolving nature of scientific knowledge.
Primordial Definitions
Historically, 'acid' originated from the Latin 'acidus', which translates to sour. This etymology reflects the sour taste often associated with acidic substances. Bases, in contrast, were identified by their characteristic slippery feel.
Arrhenius' Proposition
Svante Arrhenius, in the late 19th century, introduced a theory that:
- Acids produce hydrogen ions (H⁺) when dissolved in water. This increase in H⁺ concentration gives the solution its acidic nature.
- Bases produce hydroxide ions (OH⁻) when in an aqueous solution. The presence of these ions is responsible for the basic properties of the solution.
While groundbreaking, this definition had its limitations, mainly because it was constrained to aqueous solutions.
Image courtesy of OpenStax College
Lewis’ Electron Theory
Gilbert N. Lewis, in the early 20th century, approached the concept of acids and bases from an electronic perspective. According to his theory:
- Acids are substances that can accept an electron pair.
- Bases are substances that can donate an electron pair.
This perspective was broader and extended the definition of acids and bases beyond proton interactions, allowing chemists to explain a more diverse range of reactions.
Image courtesy of Science Notes
Brønsted–Lowry's Proton-Centric Definition
Almost concurrently with Lewis, Johannes Brønsted and Thomas Lowry put forth a similar proton-centric idea. Their definitions, as elaborated earlier, revolved around proton donation and acceptance. This theory filled the gaps left by the Arrhenius definition, offering a comprehensive explanation for reactions even outside aqueous mediums.
Wrapping up Historical Perspectives
The understanding of acids and bases has witnessed a series of refinements over time. Each subsequent theory built on its predecessor, adding depth and breadth to our comprehension. The Brønsted–Lowry definition, while building on the works of Arrhenius and others, offers a holistic and versatile understanding of acid-base interactions.
FAQ
Yes, a substance can be a Brønsted–Lowry acid but not an Arrhenius acid. An example is ammonia gas (NH3). When it reacts with hydrogen chloride gas (HCl), it forms ammonium chloride. In this reaction, ammonia accepts a proton from HCl, making it a base according to Brønsted–Lowry. However, neither NH3 nor HCl produced H+ or OH- ions in water in this reaction, so by the Arrhenius definition, neither can be classified as an acid or base. This example illustrates the broader applicability of the Brønsted–Lowry theory compared to the Arrhenius theory.
The Arrhenius theory, as previously mentioned, defines acids as substances that produce hydrogen ions in aqueous solutions and bases as substances that produce hydroxide ions. While this theory is accurate, it's restrictive in scope, only applicable to reactions in aqueous solutions. The Brønsted–Lowry theory, by focusing on proton donation and acceptance, is a more inclusive definition. It encompasses the reactions described by the Arrhenius theory (as they involve proton transfers) but also includes other proton transfer reactions outside aqueous environments. Thus, the Brønsted–Lowry theory can be seen as a broader generalisation of the Arrhenius concept.
Amphoteric substances can act as both acids and bases. This directly ties in with the Brønsted–Lowry theory, which defines acids as proton donors and bases as proton acceptors. An amphoteric substance has the flexibility to either donate or accept a proton depending on the reaction conditions and the substances it reacts with. Water, as discussed earlier, is a classic example of an amphoteric substance. The concept of amphoterism, while not exclusive to the Brønsted–Lowry theory, fits comfortably within its framework, further showcasing the theory's ability to explain a wide range of acid-base behaviours.
Water is unique in its ability to act as both an acid and a base, a property termed "amphoteric". According to the Brønsted–Lowry theory, an acid donates a proton and a base accepts a proton. In the reaction between water and ammonia, for example, water donates a proton to ammonia, acting as an acid. Conversely, when water reacts with hydrogen chloride, it accepts a proton from HCl, behaving as a base. This dual nature arises due to the molecular structure of water, allowing it to either lose a proton or gain a proton, depending on the reaction conditions.
The Brønsted–Lowry theory focuses on the transfer of protons (H+ ions) in acid-base reactions. In this theory, an acid is a proton donor and a base is a proton acceptor. The Lewis acid-base theory, on the other hand, revolves around the transfer of electron pairs. A Lewis acid is an electron-pair acceptor, while a Lewis base is an electron-pair donor. Therefore, while the Brønsted–Lowry theory restricts its definition to proton transactions, the Lewis theory is more general, encompassing a broader range of reactions through its focus on electron pairs.
Practice Questions
a) Identify the Brønsted–Lowry acid and base in the reaction.
b) Explain the difference between a base and an alkali, making sure to mention the solubility criterion.
1a: In the given reaction, NH3 accepts a proton from H2O, making it the Brønsted–Lowry base. Conversely, H2O donates a proton to NH3, so it's the Brønsted–Lowry acid.
Answer to 1b: A base, in general terms, refers to any substance that can accept a proton. It's a broad categorisation that encompasses a wide range of substances. On the other hand, an alkali is a specific kind of base that is soluble in water. This means while all alkalis are bases, not all bases are alkalis. The solubility in water is the defining criterion that differentiates alkalis from the broader category of bases.
The Arrhenius definition, proposed by Svante Arrhenius in the late 19th century, defines acids as substances that produce hydrogen ions (H+) when dissolved in water and bases as substances that produce hydroxide ions (OH-) in aqueous solutions. While this was a significant step forward, it had a significant limitation: it was constrained to reactions in aqueous solutions. This means any acid-base reactions that occurred outside of water couldn't be explained using this definition. The Brønsted–Lowry theory addressed this limitation by focusing on proton donation and acceptance, making it applicable to reactions even outside aqueous mediums.
